Iron plays an essential role in many biological processes. Complexes of iron with biomolecules are required for many vital physiologic processes, such as transport of oxygen throughout the body, synthesis of hormones, metabolism, detoxification and electron transport. It is recognized that iron deficiency can lead to many serious pathological conditions. Public health authorities recommend iron supplementation to avoid anemia caused by iron deficiency. Many forms of supplements, from elemental iron, to various salts and complexes of iron are used for oral administration.
In more severe cases of anemia, such as in patients with chronic diseases and patients undergoing dialysis, iron supplements in solution are provided by intravenous route. Many of these iron compounds have complex structures where iron atoms are held in a core surrounded by organic or inorganic moieties. Although most of these iron complexes have been known for a long time, commercial production of these compounds has been challenging and has required extensive technological developments. Typically these complexes are first made and isolated as solid and then made into a dosage form as a solution. However, the processes that involve isolation of the solid complex do not provide a product of uniform composition.
U.S. Pat. No. 7,816,404 discloses a process of preparing a water soluble ferric pyrophosphate citrate chelate by precipitating it from a liquid composition. Experimental details in Table 11 of U.S. Pat. No. 7,816,404 reveal that the solid ferric pyrophosphate citrate chelate takes about 60% by weight of the liquid composition. Extra Fe3+ may be in the form of other salts, impurities, or remains in mother liquor and thus lost in the liquid composition. The precipitated ferric pyrophosphate citrate chelate needs to be dried in an oven. Extra precautions are required during the drying to minimize decomposition of the chelate such as conversion of pyrophosphate to phosphate due to heat. It is reported that the chelate contains phosphate in an amount of 2% or less by weight.
The process disclosed in U.S. Pat. No. 7,816,404 cannot provide a product of uniform composition. Based on the data provided in Tables 10 and 11 of U.S. Pat. No. 7,816,404, we calculate percent of variations of citrate and pyrophosphate in different lots, shown as below:
Batch% Fe% Citrate% PyrophosphateNo.ScaleAssay1Assay1Variance2Assay1Variance24864 1 g8.628.1+6.2211.8−8.244865 2 g8.837.5+15.0312.7−7.814866 2 g10.232.6+6.5513.8−9.97486810 g9.630.3+5.7814.1−8.27486910 g9.727.7+3.0213.2−9.40551151610 g10.8322.4−5.1618.3−6.9451151710 g10.2723.5−2.6317.3−6.6351191325 g10.2323.4−2.6317.7−6.141Weight percents shown here are from the specific Batch No. in Tables 10 and 11 of Pat. No. 7,816,404.2Variance in citrate and pyrophosphate content is the difference in percent composition of isolated solid from the expected amount for Fe4(citrate)3(pyrophosphate)3 on the basis of assay for % Fe. It is marked as “+” if weight percent in the sample is more than expected and marked as “−” if weight percent in sample is less than expected.
The above calculations show that even when the assay for % Fe for the lots are similar, there are big differences in citrate and pyrophosphate content from lot to lot. Actually on scale up, the citrate goes from excess to deficiency, even though the same excess of citrate is used in the chemical reaction. It is obvious from these data that further scale up to commercial scale will be very challenging.
Without wishing to be bound by theory, the inconsistency in net output with scale given same input, as observed in the prior art process, may be caused by loss on drying and degradation (e.g., from pyrophosphate to phosphate) due to heat, as required by the process of U.S. Pat. No. 7,816,404.
U.S. Pat. No. 8,178,709 reports a method of preparation of a water soluble ferric citrate chelate with varying amounts of pyrophosphate and sodium, which indicates that non-essential components make up a significant weight of the bulk mass. This material is prepared by air oxidation of Fe2+ to Fe3+ and requires 3-7 days to complete the reaction in laboratory. The composition that enables the oxidation is essentially free of sulfate. Because the product is light sensitive, the oxidation process has to be run in dark. Given the long reaction time, scaling up the oxidation process for commercial production is very challenging and requires many operational controls.
Therefore, there is a need for an operationally simple procedure for preparing solutions of iron salt that can readily be administered.